1. University of Louisiana at Lafayette
Chapter 1
Lecture
Outline
Prepared by
Andrea D. Leonard
University of Louisiana at Lafayette
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. 1.1 Chemistry—The Science of Everyday Experience
Chemistry is the study of matter—its composition,
properties, and transformations.
Matter is anything that has mass and takes up volume.
Matter can be:
Naturally occurring:
cotton
sand
digoxin, a cardiac drug
Synthetic (human-made):
nylon
Styrofoam
ibuprofen

3. 1.2 States of Matter The Solid State: A solid has a definite volume.
It maintains its shape regardless of its container.
Solid particles lie close together in a regular pattern.

4. 1.2 States of Matter The Liquid State: A liquid has a definite volume.
It takes the shape of its container.
Liquid particles are close together but can move past one another.

5. 1.2 States of Matter The Gas State:
A gas has no definite shape; it assumes the shape of its container.
It has no definite volume; it assumes the volume of its container.
Gas particles are very far apart and move around randomly.

6. 1.2 States of Matter Physical properties can be observed or measured
without changing the composition of the material.
boiling point
melting point
solubility
color
odor
state of matter

7. 1.2 States of Matter
A physical change alters the material without changing
its composition (changes in state).

8. 1.2 States of Matter Chemical properties determine how a substance can
be converted into another substance.
Chemical change is the chemical reaction that
converts one substance into another (Chapters 5 and 6).

9. 1.3 Classification of Matter
All matter can be classified as either a pure substance or a mixture.
I. Pure Substances
A pure substance is composed of only a single component (atom or molecule).
It has a constant composition, regardless of sample size or origin of sample.
It cannot be broken down to other pure substances by a physical change.

10. 1.3 Classification of Matter
All matter can be classified as either a pure substance or a mixture.
I. Pure Substances
Table sugar (C12H22O11) and water (H2O) are both
pure substances:

11. 1.3 Classification of Matter
All matter can be classified as either a pure substance or a mixture.
II. Mixtures
Mixtures are composed of more than one component.
They can have varying composition (any combination of solid, liquid, and gas).
Mixtures can be separated into their components by a physical process.

12. 1.3 Classification of Matter
All matter can be classified as either a pure substance or a mixture.
II. Mixtures
Sugar dissolved in water is a mixture.

13. 1.3 Classification of Matter
A pure substance is classified as an element or a compound.
I. An element is a pure substance that cannot be broken down by a chemical change.
aluminum metal (Al)

14. 1.3 Classification of Matter
A pure substance is classified as an element or a compound.
II. A compound is a pure substance formed by chemically joining two or more elements.
table salt (NaCl)

15. 1.3 Classification of Matter

16. 1.4 Measurement Every measurement is composed of a number and a unit.
The number is meaningless without the unit.
Examples:
proper aspirin dosage = 325 (milligrams or pounds?)
a fast time for the 100-meter dash = (seconds or days?)

17. 1.4 Measurement A. The Metric System
Each type of measurement has a base unit in the metric system..

18. 1.4 Measurement A. The Metric System
Other units are related to the base unit by a power of 10.
The prefix of the unit name indicates if the unit is larger or smaller than the base unit.

19. 1.4 Measurement B. Measuring Length
The base unit of length is the meter (m).
1 kilometer (km) = 1,000 meters (m)
1 km = 1,000 m
1 millimeter (mm) = meters (m)
1 mm = m
1 centimeter (cm) = 0.01 meters (m)
1 cm = 0.01 m

20. 1.4 Measurement C. Measuring Mass
Mass is a measure of the amount of matter in an object.
Weight is the force that matter feels due to gravity.
The base unit of mass is the gram (g).
1 kilogram (kg) = 1,000 grams (g)
1 kg = 1,000 g
1 milligram (mg) = grams (g)
1 mg = g

21. 1.4 Measurement D. Measuring Volume
The base unit of volume is the liter (L).
1 kiloliter (kL) = 1,000 liters (L)
1 kL = 1,000 L
1 milliliter (mL) = liters (L)
1 mL = L
Volume = Length x Width x Height
= cm x cm x cm
= cm3
1 mL = 1 cm3 = 1 cc

22. 1.4 Measurement

23. 1.5 Significant Figures
An exact number results from counting objects or is
part of a definition.
10 fingers
10 toes
1 meter = 100 centimeters
An inexact number results from a measurement or
observation and contains some uncertainty.
15.3 cm g mL

24. 1.5 Significant Figures A. Determining Significant Figures
Significant figures are all the digits in a measured
number including one estimated digit.
All nonzero digits are always significant.
65.2 g
65.2 g g g
3 sig. figures
6 sig. figures

25. 1.5 Significant Figures A. Determining Significant Figures
Rules for Zero:
Rule 1: A zero counts as a significant figure when
it occurs:
between two nonzero digits
29.05 g
29.05 g mL mL
4 sig. figures
5 sig. figures
at the end of a number with a decimal place cm cm
620. lb
620. lb
5 sig. figures
3 sig. figures

26. 1.5 Significant Figures A. Determining Significant Figures
Rules for Zero:
Rule 2: A zero does not count as a significant figure
when it occurs:
at the beginning of a number mg mg
0.008 mL
0.008 mL
3 sig. figures
1 sig. figure
at the end of a number that does not have a decimal
2570 m
2570 m m m
3 sig. figures
5 sig. figures

27. 1.5 Significant Figures B. Rules for Multiplication and Division
The answer has the same number of significant figures
as the original number with the fewest significant figures.
4 sig. figures
351.2 miles
351.2 miles
miles =
hour
5.5 hour
5.5 hour
2 sig. figures
Answer must have
2 sig. figures.

28. 1.5 Significant Figures B. Rules for Multiplication and Division
to be retained
to be dropped
miles =
64 miles
hour
hour
first digit to be dropped
2 sig. figures
Answer
If the first digit
to be dropped is:
Then:
between 0 and 4
drop it and all remaining digits
between 5 and 9
round up the last digit
to be retained by adding 1

29. 1.5 Significant Figures B. Rules for Multiplication and Division

30. 1.5 Significant Figures C. Rules for Addition and Subtraction
The answer has the same number of decimal places
as the original number with the fewest decimal places.
10.11 kg
10.11 kg
2 decimal places
3.6 kg
3.6 kg
1 decimal place
6.51 kg
answer must have
1 decimal place
final answer
1 decimal place =
6.5 kg

31. In scientific notation, a number is written as:
y x 10x
y x 10x
Exponent:
Any positive or negative
whole number.
Coefficient:
A number between
1 and 10.

32. 1.6 Scientific Notation 2,500 0.036 2500 0.036 2.5 x 103 3.6 x 10−2
HOW TO Convert a Standard Number to Scientific Notation
Example
Convert these numbers to scientific notation.
2,500
0.036
Move the decimal point to give a number
between 1 and 10.
Step [1]
2500
0.036
Multiply the result by 10x, where
x = number of places the decimal was moved.
Step [2]
move decimal left,
x is positive
move decimal right,
x is negative
2.5 x 103
3.6 x 10−2

33. Converting a Number in Scientific Notation
to a Standard Number
When the exponent x is positive, move the decimal point x places to the right.
2.800 x 102 =
280.0
When the exponent x is negative, move the decimal point x places to the left.
2.80 x 10–2 =
0.0280

34. 1.7 Using the Factor-Label Method A. Conversion Factors
Conversion factor: A term that converts a quantity in
one unit to a quantity in another unit.
original
quantity
conversion factor
desired x =
Conversion factors are usually written as equalities.
2.21 lb = 1 kg
To use them, they must be written as fractions.
2.21 lb
1 kg or

35. units are treated like numbers make sure all unwanted units cancel
1.7 Using the Factor-Label Method B. Solving a Problem Using One Conversion Factor
Factor-label method: Using conversion factors to convert a quantity in one unit to a quantity in
another unit.
units are treated like numbers
make sure all unwanted units cancel
To convert 130 lb into kilograms:
130 lb x
conversion factor
? kg =
original quantity
desired quantity

36. 1. 7 Using the Factor-Label Method B
1.7 Using the Factor-Label Method B. Solving a Problem Using One Conversion Factor
2.21 lb
1 kg
Answer
2 sig. figures
130 lb x or
1 kg
2.21 lb =
59 kg
The bottom conversion factor has
the original unit in the denominator.
The unwanted unit lb cancels.
The desired unit kg does not cancel.

37. 1.7 Using the Factor-Label Method
HOW TO Solve a Problem Using Conversion Factors
How many grams of aspirin are in a 325-mg
tablet?
Example
Identify the original quantity and the desired
quantity, including units.
Step [1]
original quantity
desired quantity
325 mg
? g

38. 1.7 Using the Factor-Label Method
HOW TO Solve a Problem Using Conversion Factors
Write out the conversion factor(s) needed
to solve the problem.
Step [2]
1 g = 1000 mg
This can be written as two possible fractions:
1000 mg 1g
1 g
1000 mg or
Choose this factor to
cancel the unwanted
unit, mg.

39. 1.7 Using the Factor-Label Method
HOW TO Solve a Problem Using Conversion Factors
Step [3]
Set up and solve the problem.
1 g
1000 mg
0.325 g
0.325 g
325 mg
325 mg x =
3 sig. figures
3 sig. figures
Unwanted unit
cancels
Write the answer with the correct number
of significant figures.
Step [4]

40. 1. 7 Using the Factor-Label Method C
1.7 Using the Factor-Label Method C. Solving a Problem Using Two or More Conversion Factors
Always arrange the factors so that the denominator in
one term cancels the numerator in the preceding term.
How many liters is in 1.0 pint?
1.0 pint
original quantity
? L
desired quantity
Two conversion factors are needed:
2 pints = 1 quart
1.06 quarts = 1 liter
2 pt
1 qt
1 qt
2 pt
1.06 qt
1 L
1 L
1.06 qt or or
First, cancel pt.
Then, cancel qt.

41. Set up the problem and solve:
1.7 Using the Factor-Label Method C. Solving a Problem Using Two or More Conversion Factors
Set up the problem and solve:
1 qt
2 pt
1 L
1.06 qt
1.0 pt
1.0 pt x x =
0.47 L L
2 sig. figures
2 sig. figures
Write the answer with the correct number of
significant figures.

42. To convert from oC to oF: To convert from oF to oC:
1.9 Temperature
Temperature is a measure of how hot or cold an object is.
Three temperature scales are used:
Degrees Fahrenheit (oF)
Degrees Celsius (oC)
Kelvin (K)
To convert from oC to oF:
To convert from oF to oC:
oC = oF − 32
1.8
oF = 1.8(oC)
To convert from oC to K:
To convert from K to oC:
K = oC
oC = K − 273

43. 1.9 Temperature Comparing the Three Temperature Scales

44. 1.10 Density and Specific Gravity A. Density
Density: A physical property that relates the mass of
a substance to its volume.
mass (g)
density =
volume (mL or cc)
To convert volume (mL) to mass (g):
To convert mass (g) to volume (mL): g mL mL x = g g x = mL mL g
inverse of density
density

45. 1.10 Density and Specific Gravity A. Density
Example:
If the density of acetic acid is 1.05 g/mL, what is
the volume of 5.0 grams of acetic acid?
5.0 g
? mL
original quantity
desired quantity
Density is the conversion factor, and can be
written two ways:
1.05 g
1 mL
1 mL
1.05 g
Choose the inverse density
to cancel the unwanted unit, g.

46. 1.10 Density and Specific Gravity A. Density
Set up and solve the problem:
1 mL
1.05 g
5.0 g
5.0 g x = mL
4.8 mL
2 sig. figures
2 sig. figures
Unwanted unit
cancels
Write the final answer with the correct number
of significant figures.

47. 1.10 Density and Specific Gravity B. Specific Gravity
Specific gravity: A quantity that compares the density
of a substance with the density of water at the
same temperature.
density of a substance (g/mL)
density of water (g/mL)
specific gravity =
The units of the numerator (g/mL) cancel the
units of the denominator (g/mL).
The specific gravity of a substance is equal to its
density, but contains no units.